QD 


Ft-Ws 


DflD 


EXCHANGE 


Oxidation  and  Reduction  without 
the  Addition  of  Acid 

I.     The  Reaction  between  Ferrous  Sulfate  and 
Potassium  Dichromate 

II.     The  Reaction  between  Stannous  Chloride 
and  Potassium  Dichromate 


A  DISSERTATION 

SUBMITTED  TO    THE    FACULTY  OF  THE  GRADUATE  SCHOOL   OF    TH 
IXIYERSITY  OF  PITTSBURGH  IN  CONFORMITY  WITH  THE 
REQUIREMENTS  FOR  THE  DEGREE  OF 
DOCTOR  OF  PHILOSOPHY 


BY 


JOSHUA  CHITWOOD  WITT 


PITTSBURGH 
1915 


EASTON,  PA.: 

ESCHENBACH  PRINTING  Co. 
1916 


Oxidation  and  Reduction  without 
the  Addition  of  Acid 

I.     The  Reaction  between  Ferrous  Sulfate  and 
Potassium  Bichromate 

II.     The  Reaction  between  Stannous  Chloride 
and  Potassium  Dichromate 


A  DISSERTATION 

SUBMITTED  TO    THE    FACULTY  OF  THE  GRADUATE  SCHOOL   OF    THE 

UNIVERSITY  OF  PITTSBURGH  IN  CONFORMITY  WITH  THE 

REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


BY 

JOSHUA  CHITWOOD  WITT 

PITTSBURGH 
1915 


EASTON,  PA.: 

ESCHENBACH  PRINTING  Co. 
1916 


ACKNOWLEDGMENT. 

The  writer  is  greatly  indebted  to  Dr.  Marks  Neidle,  under  whose  super- 
vision this  work  has  been  carried  out,  for  his  kind  consideration  and  help- 
fulness. 


OXIDATION  AND   REDUCTION  WITHOUT  THE  ADDITION 

OF  ACID. 

I.    THE  REACTION   BETWEEN   FERROUS   SULFATE   AND   POTASSIUM 

BICHROMATE. 
BY  JOSHUA  C.  WITT. 

The  first  use  of  the  reaction  between  ferrous  salts  and  dichromate 
for  the  determination  of  iron  was  made  by  Penny.1  In  the  method, 
as  described  in  his  paper,  a  sample  of  "iron  stone"  was  dissolved  in  hydro- 
chloric acid,  and  the  iron  reduced  by  adding  sodium  sulfite  in  excess.  After 
boiling  off  the  excess  sulfurous  acid,  he  titrated  with  dichromate  solution, 
using  potassium  ferricyanide  as  an  outside  indicator.  It  is  interesting 
to  note  that  this  method  is  essentially  the  same  as  that  in  use  today  for 
the  determination  of  iron  in  iron  ore. 

Whenever  the  reaction  between  ferrous  salts  and  dichromate  has  been 
studied  a  mineral  acid  has  been  added.  Penny  employed  excess  of  free 
acid  in  dissolving  the  iron  ore,  and  the  equation  for  the  reaction  demands 
free  acid  for  the  formation  of  the  normal  salts  of  potassium,  chromium, 
and  iron.  Since  no  mention  of  any  investigation  of  the  reaction  in  the 
absence  of  free  acid  could  be  found  in  the  literature,  it  was  decided  to 
perform  a  few  preliminary  experiments  in  which  a  quantity  of  ferrous 
sulfate  was  titrated  by  o.i  N  dichromate,  with  and  without  acid. 

It  was  considered  preferable  to  weigh  out  a  separate  portion  of  ferrous 
sulfate  for  each  titration,  rather  than  to  keep  a  standard  solution  of  the 
salt.  As  soon  as  a  portion  was  weighed  out  it  was  rapidly  transferred 
to  a  beaker  containing  water,  and  titrated  at  once  with  the  dichromate 
solution,  using  potassium  ferricyanide  as  an  outside  indicator.  The 
following  results  were  obtained  showing  the  effect  of  acid: 

FeSO4.7H»O  (g.).  Cc.  KjCrjOr.  HiSO4. 

0.8  29.52  Excess  present 

0.8  29.54  Excess  present 

0.8  36.16  None  present 

0.8  36.25  None  present 

The  end  point  is  obtained  when  the  amount  of  ferrous  salt  remaining 
at  the  time  the  drop  test  is  made  is  insufficient  to  affect  the  indicator. 
When  no  acid  is  added,  an  excess  of  dichromate  is  required  to  give  an  end 
point,  which  means  that  with  the  theoretical  amount  of  dichromate 
necessary  to  completely  oxidize  the  ferrous  sulfate,  enough  of  the  latter 
remains  to  affect  the  indicator,  i.  e.,  the  reaction  is  incomplete.  A  brown 
precipitate  appears  after  a  few  drops  of  the  dichromate  have  been  added. 

The  following  results  show  the  effect  of  the  volume  of  ferrous  sulfate 
solution  on  the  titration,  0.8  g.  of  salt  being  used  in  each  experiment, 
1  Brit.  Assoc.  Rep.,  [2]  1850,  58,  59. 


Cc.  water  Cc.  dichromate  Cc.  dichromate 

added  to  PeSO«.  sol.  with  HjSO*.  sol.  without  HjS 

o  29.67  29.89 

5           :....  29.88 

15  30.03 

30  30.62 

IOO  32  .  12 

looo  53-00 

This  increase  in  the  dichromate  was  to  be  expected,  since  the  reaction 
is  slower  the  greater  the  volume,  and  larger  amounts  of  dichromate  are 
required  to  drive  the  reaction  to  the  end  point.  When  no  water  is  added 
the  result  of  the  titration  is  nearer  theoretical,  and  in  several  experiments, 
in  which  more  than  0.8  g.  was  taken  and  the  solid  titrated,  the  result  was 
exactly  the  theoretical.  We  may  therefore  conclude  that  the  precipitate 
formed  does  not  adsorb  the  ferrous  ion  appreciably.  Adsorption  of 
ferrous  ion  would  vitiate  the  results  on  the  velocity  of  the  reaction. 

Considerable  difficulty  was  encountered  in  finding  the  end  point  at 
the  higher  concentrations  when  the  titration  was  made  in  the  absence 
of  sulfuric  acid.  The  brown  precipitate  had  a  tendency  to  mask  the 
end  point.  To  overcome  this,  when  the  end  point  was  nearly  reached, 
it  was  found  necessary  to  filter  a  few  drops  of  the  mixture  each  time  be- 
fore it  was  applied  to  the  indicator. 

Measurement  of  the  Velocity  of  the  Reaction. — The  problem  which 
presented  itself  at  this  point  was  to  find  a  method  of  determining  the 
unoxidized  ferrous  salt  or  unreduced  dichromate  in  a  solution  containing 
ferric  salts,  ferrous  salts,  chromic  salts,  and  dichromate.  Three  methods 
suggested  themselves: 

(1)  To  stop  the  reaction  by  the  addition  of  ammonium  hydroxide, 
filter  the  precipitated  hydroxides  of  iron  and  chromium,  and  determine 
the  chromium  in  the  precipitate. 

(2)  To  add  ammonium  hydroxide  as  in  (i)  and  titrate  the  unchanged 
dichromate  in  the  filtrate. 

(3)  To  precipitate  the  unchanged  dichromate  with  lead  acetate,  dis- 
solve the  precipitate  of  the  reaction  in  acetic  acid,  and  determine  chro- 
mate  in  the  residue. 

The  second  method,  being  more  direct  and  therefore  more  accurate, 
was  adopted. 

It  is  well  known  that  ferrous  salts,  in  common  with  salts  of  other  di- 
valent metals,  cannot  be  completely  precipitated  by  ammonium  hydroxide 
in  the  presence  of  ammonium  salts  in  consequence  of  the  repression  of 
hydroxyl  ion  by  the  latter.  In  order  to  be  certain  of  completely  precipi- 
tating ferrous  iron,  the  necessary  conditions  were  investigated. 

It  was  found  that  ferrous  salts  may  be  completely  precipitated  with 
ammonium  hydroxide  provided, 


(1)  The  solution  is  neutral. 

(2)  No  ammonium  salts  are  present  to  begin  with. 

(3)  The  concentration  is  sufficiently  low. 

(4)  The  solution  is  boiled  and  the  precipitate  allowed  to  settle  before 
filtration  is  attempted. 

If  0.5  g.  portions  of  ferrous  sulfate  were  dissolved  in  various  volumes 
of  water,  and  an  excess  of  ammonium  hydroxide  added,  the  precipitation 
was  complete  only  when  the  volume  was  at  least  100  cc. 

Solutions. 

Potassium  Dichromate. — A  o.i  N  solution,  standardized  against  iron 
wire,  was  kept  in  a  ten-liter  bottle  fitted  with  a  siphon.  All  air  entering 
the  bottle  came  through  a  cotton  plug  to  avoid  contamination. 

Sodium  Thiosulfate. — A  o.oi  N  solution  was  standardized  each  time 
before  using  against  the  dichromate  solution. 

Ferrous  Sulfate. — At  first  it  was  thought  advisable  to  make  up  a  stand- 
ard solution  of  ferrous  sulfate  and  attempt  to  protect  it  from  oxidation, 
but  it  was  finally  decided  to  use  the  dry  salt  and  weigh  out  a  portion  for 
each  determination.  To  avoid  difficulty  from  any  variation  in  quality, 
a  fresh  pound  bottle  of  the  c.  P.  salt  was  taken,  and  used  for  all  the  work. 
The  surface  layer  was  discarded,  a  weighing  bottle  filled  and  kept  in  the 
balance  until  used,  then  refilled  when  necessary.  In  weighing  out  a  sample, 
a  slight  excess  was  placed  on  the  balance  and  the  stopper  of  the  weighing 
bottle  replaced  at  once.  The  excess  salt  was  removed  as  quickly  as 
possible  and  discarded  to  avoid  any  possible  contamination.  The  salt 
was  analyzed  from  time  to  time  and  found  to  remain  constant  in  compo- 
sition, as  shown  by  the  following  results  obtained  with  0.8  g.  samples: 

Date April  5.  May  3.  June  9. 

Titration  with  0.1014  N  K2Cr2O7 29.53  cc.  29.67  cc.  29.63  cc. 

Manipulation. — A  large,  electrically  controlled  bath  was  maintained 
at  30°  =±=  0.05°.  A  ten-liter  bottle  of  distilled  water  was  kept  in  this 
bath  that  no  delay  might  be  caused  by  waiting  for  water  to  assume  the 
correct  temperature. 

Nearly  as  much  water  as  was  needed  for  the  experiment  was  placed  in 
a  liter  flask  corrected  for  temperature,  and  a  given  amount  of  standard 
dichromate  solution  was  run  into  an  Erlenmeyer  flask.  Both  flasks 
were  immersed  in  the  bath  and  allowed  to  assume  constant  temperature. 
In  the  meantime  a  portion  of  the  ferrous  sulfate  was  weighed  and  rapidly 
transferred  to  the  liter  flask.  When  solution  was  complete,  the  dichromate 
was  added  and  the  volume  adjusted.  The  flask  was  kept  in  the  bath 
and,  at  intervals,  100  cc.  portions  were  removed  and  run  into  beakers 
containing  excess  of  ammonium  hydroxide. 

The  precipitate  formed  by  the  ammonium  hydroxide  was  very 
finely  divided  and  would  pass  very  readily  through  the  filter.  It  was 


8 

found  necessary  to  let  it  stand  some  time — preferably  over  night — or  to 
boil  it  a  few  minutes  before  a  complete  filtration  could  be  made.  It  was 
preferable  to  filter  at  once,  without  heating,  but  no  method  could  be 
found  which  gave  the  desired  result.  The  precipitate  passed  through  an 
alundum  Gooch,  and  would  not  settle  when  kept  in  a  centrifuge  for  15-20 
minutes. 

In  order  to  determine  whether  the  potassium  dichromate  still  in  solution 
was  in  any  way  affected  by  the  precipitated  ferrous  iron,  25  cc.  of  o.i  N 
dichromate  was  added  to  0.8  g.  of  ferrous  sulfate  dissolved  in  water  in  a 
liter  flask,  and,  after  introducing  an  excess  of  ammonium  hydroxide, 
the  mixture  was  made  up  to  volume.  A  number  of  100  cc.  portions  were 
withdrawn  and  placed  in  beakers.  They  were  filtered  at  various  intervals 
and  titrated  with  o.oi  TV  thiosulfate  by  the  method  already  described. 
Some  were  boiled  before  filtering,  and  others  were  filtered  in  the  cold. 
The  following  are  the  results  obtained : 

Time.  Titration,  0.01  N  thiosulfate.  Remarks. 

2  hours  o .  90  cc.  Not  boiled 

2  hours  i .  oo  Boiled 

24  hours  o .  94  Not  boiled 

24  hours  i .  oo  Boiled 

It  is  seen  from  the  above  that  the  final  result  is  not  altered  by  allowing 
the  mixture  to  stand  for  many  hours,  or  by  boiling,  before  filtration. 

Results. — All  measurements  were  made  at  30°.  The  ferrous  sulfate 
and  dichromate  solution  were  in  the  ratio  of  0.8  g.  of  the  former  to  25  cc. 
of  the  latter,  or  2.878  mols  to  0.4225  mol.  It  was  not  thought  advisable 
to  attempt  any  measurements  with  more  dichromate  than  would  be  re- 
quired for  the  normal  end  point,  since  in  this  case  a  very  large  volume  of 
o'.oi  N  thiosulfate  would  be  required.  The  ratio  of  the  reacting  substances 
was  maintained  constant.  The  volume  of  dichromate  reduced  is  repre- 
sented by  x  and  that  unreduced  by  a  —  x. 

TABLE  I. 
Total  Volume  Containing  25  cc.  of  0.1014  N  K2Cr2O7  and  0.8  g.  FeSO4. 

100  cc.  250  cc.  500  cc.  1,000  cc.          2,000  cc.          4,000  cc.         5,000  cc. 

utes.  a  —  x.      x.     a  —  *.       x.     a  —  x.       x.     a  —  *.      x.     a  —  x.      x.  a  —  x.     x.      a  —  x.       x. 

I     0.06   24.94  0.22   24.78  0.35   24.65    1.50  23.50   1.93   23.07  

5     0.03  24.97  0.18  24.82  0.34  24.66   1.46  23.50    1.99  23.01   2.69  23.31 

15      0.04    24.96    0.23    24.77    0.30    24.70   0-93    24.07    1.74    23.26    1.65    23.35    2.20    23.80 
30     0.04    24.96    0.22    24.78    0.29    24.71    0.80    24.20    1.36    23.64    1.57    23.43    2.20    23.80 

60    0.04  24.96  0.19  24.81  0.26  24.74  0.61   24.37    i .06  23.94  2.31   23.69 

One  series  of  experiments  was  made  with  method  number  three  as  a 
check.  The  work  was  carried  on  in  the  same  way  up  to  the  time  when  100 
cc.  portions  were  removed  from  the  liter  flask.  In  this  case  they  were 
run  into  beakers  containing  lead  acetate  solution,  which  precipitated  the 


sulfate  ion  and  the  chromate  ion.  The  mixture  was  then  acidified  with 
a  few  cubic  centimeters  of  acetic  acid  and  boiled  to  dissolve  all  the  iron 
salts.  The  lead  salts  were  then  filtered  out  and  the  lead  chromate  dis- 
solved in  dilute  hydrochloric  acid.  The  resulting  dichromate  was  titrated, 
after  cooling,  with  o.oi  N  thiosulfate.  The  results  given  in  Table  II 
compare  satisfactorily  with  those  previously  obtained  and  given  in 
Table  I. 

TABLE  II. 


Minutes.  Method  III.  Method  II. 

5  1-49  1.46 

15  0-98  0.93 

60  0.59  0.6i 

Comments  on  Velocity  Measurements  and  the  Order  of  the  Reaction.— 
From  Table  I  it  is  seen  that  in  the  titration  of  0.8  g.  of  ferrous  sulfate 
with  dichromate,  the  reaction  is  99.8%  complete  at  the  end  of  one  minute, 
provided  the  final  volume  is  100  cc.  This  statement  may  be  made  even 
though  in  our  experiments  the  dichromate  taken  was  a  little  less  than 
equivalent  to  the  ferrous  sulfate.  At  all  other  concentrations  except 
the  most  dilute,  the  reaction  is  more  than  90%  complete  at  the  end  of 
the  first  minute. 

The  data  as  obtained  are  not  of  a  nature  to  permit  ready  calculation 
of  the  order  of  the  reaction,  although  those  in  the  last  column  of  Table 
I  seemed  sufficiently  regular  to  justify  an  attempt  at  such  a  calculation. 
No  constant  could  be  obtained  by  assuming  the  reaction  to  be  of  the 
first  order  with  respect  to  each  of  the  reacting  substances,  of  the  first  order 
with  respect  to  one  and  of  the  second  with  respect  to  the  other,  and,  finally, 
of  the  second  order  with  respect  to  both.  Our  conclusion,  therefore,  is 
that  this  reaction  is  probably  of  an  order  higher  than  the  fourth. 

The  rate  of  oxidation  of  ferrous  sulfate  by  dichromate  with  the  addition 
of  more  than  the  sulfuric  acid  required  by  the  normal  equation  has  been 
investigated  by  Benson.1  It  is  stated  in  his  conclusions  that  the  rate  is 
proportional  to  the  second  power  of  the  concentration  of  ferrous  salt, 
and  to  the  second  power  of  that  of  the  acid,  and  that  the  order  is  variable 
with  respect  to  the  dichromate.  Benson  also  found  that  the  order  is 
much  retarded  by  the  presence  of  ferric  salts.  If  the  velocity  of  this  re- 
action is  strictly  proportional  to  the  square,  or  any  other  power,  of  the 
concentration  of  acid  added  it  should  be  zero  when  no  acid  is  employed. 

There  can  be  no  question,  however,  that  the  velocity  is  proportional 
to  some  power  of  the  hydrogen-ion  concentration,  in  which  case  the  ve- 
locity of  the  reaction  without  the  addition  of  acid  is  due  to  the  hydrogen- 
ion  concentration  arising  from  the  hydrolytic  dissociation  of  both  dichro- 
mate and  ferrous  salt.  The  concentration  of  hydrogen  ion  must  play  an 
1  /.  Phys.  Chem.,  i,  i  (1903). 


10 

important  part  in  the  reaction,  even  in  very  low  concentration.  Our 
reaction  is  most  probably  accompanied  by  a  change  in  the  concentration 
of  hydrogen  ion,  which  was  disregarded  in  our  velocity  calculations.  For 
this  reason,  we  can  not  conclude  with  certainty  that  the  reaction  is  of  an 
order  higher  than  the  fourth. 

The  great  velocity  of  the  reaction  without  the  addition  of  acid  is  partly 
due  to  the  fact  that  less  than  one-third  of  the  iron  remains  in  solution 
as  ferric  salt,  which  has  a  retarding  influence,  while  the  remainder  pre- 
cipitates in  the  form  of  hydrous  ferric  oxide  and  adsorbed  ferric  sulfate. 

The  Products  of  the  Reaction. — Preliminary  experiments  showed  that 
all  the  brown  precipitate,  ultimately  formed  when  solutions  of  dichromate 
and  ferrous  sulfate  are  mixed,  does  not  come  down  instantly,but  gradually, 
reminding  one  of  the  precipitation  of  suspensoids  by  small  quantities  of 
electrolytes.  Upon  filtering  the  mixture  after  it  had  stood  for  several 
days,  the  filtrate  still  yielded  apparently  the  same  precipitate  on  standing. 
The  precipitation,  it  was  found,  could  be  rendered  complete  by  boiling, 
when  a  reddish  brown,  gelatinous  precipitate,  resembling  ferric  hydroxide, 
appeared. 

One-tenth  of  the  equivalent  weights  of  potassium  dichromate  and  ferrous 
sulfate  were  dissolved  in  water  and  the  solutions  mixed,  diluted  nearly 
to  a  liter  and  heated  to  boiling  for  several  minutes  to  bring  about  com- 
plete precipitation.  After  cooling,  the  mixture  was  made  up  to  a  liter 
exactly.  The  precipitate  was  brown  and  very  abundant.  The  superna- 
tant liquid  had  the  purplish  green  color  characteristic  of  chromium  salts. 

It  was  thought  that  heating  the  mixture  might  have  some  effect  on 
the  reaction.  To  settle  this  point,  another  solution  was  made  up  exactly 
like  the  one  already  described,  except  that  it  was  not  heated.  After 
standing  over  night,  it  was  filtered  and  both  precipitate  and  filtrate  were 
analyzed  along  with  those  of  the  first  mixture,  giving  practically  the 
same  results.  Although  the  precipitation  was  not  complete,  the  differ- 
ence was  practically  negligible.  The  work  on  this  second  solution  was 
dropped,  therefore,  and  only  the  first  carried  on. 

The  brown  precipitate  from  the  first  mixture  was  dried  to  constant 
weight  at  100-105°,  giving  a  very  hygroscopic  amorphous  powder.  This 
solid  and  also  the  filtrate  were  analyzed  for  SO3,  Cr2O3,  and  Fe2O3. 

The  SOs  was  determined  as  BaSO4.  To  determine  iron  and  chromium, 
the  hydrochloric  acid  solution  was  neutralized  with  sodium  hydroxide 
and  the  chromium  oxidized  by  sodium  peroxide.  After  boiling,  the  ferric 
hydroxide  was  filtered  out  and  washed  with  hot  water.  The  precipitate 
was  then  dissolved  in  hot  hydrochloric  acid,  reprecipitated  with  sodium 
hydroxide,  again  treated  with  sodium  peroxide,  filtered  and  washed. 
The  two  filtrates  were  combined,  boiled,  acidified  with  hydrochloric  acid 
(5  cc.  in  excess)  and  again  boiled  for  some  time.  After  cooling,  10  cc.  of 


Grams. 

Gram                              Grams 
equivalents.             originally  present. 

SOs.... 
Fe  

.    6.241 
,     1.176 
1.  122 

0.1560 
0.0631 
o  .  0647 

PRECIPITATE. 

8.006 
5-590 
1-733 

Cr  

Fe2O3 

Grams. 
6.313 

Gram  equivalents. 
0.2369 
0.0353 

o  .  0440 
o.oon 

Cr2O3 

o  894. 

SOs 

I    7SS 

Loss  on  ignitic 
Undetermined 

»n  (except  SO3)  .    . 

2  .  239 

(K20).. 

.    0.067 

II 

a  10%  potassium  iodide  solution  was  added,  and  the  solution  titrated 
with  o.  i  N  thiosulfate,  using  starch  as  the  indicator. 

The  iron  was  again  dissolved,  brought  nearly  to  dryness  on  the  hot 
plate,  reduced  by  stannous  chloride  and  titrated  with  o.i  N  potassium 
permanganate. 

FILTRATE. 

Percentage  of  total 
in  precipitate. 

22.05 
78.96 
35-26 

Percentage. 

56.03 

7-93 

15-58 
19.87 
0.59 

Further  Investigation  of  the  Precipitate. — It  is  seen  that  the  pre- 
cipitate contains  quantities  of  all  the  salts  produced  in  the  reaction. 
In  order  to  ascertain  the  nature  of  these  adsorbed  salts,  a  weighed  portion 
of  the  precipitate  was  boiled  in  water  for  some  minutes,  and  filtered. 
The  filtrate  was  made  up  to  250  cc.,  and  25  cc.  portions  removed  for  anal- 
ysis. It  was  found  that  the  nitrate  contained  2. 1 1%  SO3,  calculated  on  the 
basis  of  the  amount  of  precipitate  taken;  or,  13.54%  of  the  SO3  present 
in  the  original  precipitate  had  been  removed  by  the  first  boiling. 

A  50  cc.  portion  of  the  filtrate,  analyzed  for  iron,  gave  0.69%,  approxi- 
mately the  amount  required  to  correspond  to  the  formula  Fe2(SO4)3. 
We  may  therefore  conclude  that  the  adsorbed  salt  is  mainly  Fe2(SO4)3. 

Discussion  of  Results  on  the  Products  of  the  Reaction. — The  value 
35.26%  for  the  amount  of  chromium  in  the  precipitate  suggests  that 
one-third  of  the  total  is  precipitated  as  Cr2O3  and  the  remaining  1.93% 
adsorbed  as  Cr2(SO4)3.  If  we  add  the  number  of  equivalents  corresponding 
to  the  adsorbed  potassium  sulfate  and  chromium  sulfate,  and  subtract  the 
sum  from  the  total  number  of  equivalents  of  SO3  in  the  precipitate,  the 
result  gives  the  number  of  equivalents  of  Fe2(SO4)3  adsorbed.  This  value 
is  0.0407,  which,  added  to  the  number  of  equivalents  of  Fe2(SO4)3  in  the 
filtrate  (0.0631),  gives  the  number  of  equivalents  of  this  salt  formed  in  the 
reaction  (0.1038).  The  mixture  contained  sufficient  iron  for  0.3  equiva- 
lents of  Fe2(SO4)3.  Thus  two-thirds  of  the  iron  forms  hydrous  ferric 
oxide,  and  one- third  forms  ferric  sulfate. 

The  following  equation  completely  harmonizes  with  the  above  results: 
3K2Cr207  +  i8FeS04  +  (x  +  6y)H2O  = 

Cr2O3.*H20  +  2Cr2(SO4)3  -f  3Fe2(SO4)3 


12 

where  Cr2O3.#H2O  and  Fe2O3.;yH2O  stand  for  the  colloidal  oxides  of  chro- 
mium and  iron,  each  carrying  adsorbed  water. 

The  products  of  the  reaction  between  potassium  dichromate  and  ferrous 
sulfate  without  the  addition  of  acid  are:  potassium  sulfate,  chromium 
sulf ate  and  colloidal  chromic  oxide  in  the  molar  ratio  of  2:1;  and  ferric 
sulfate  and  colloidal  ferric  oxide  in  the  molar  ratio  of  1:2.  The  colloids 
are  precipitated  by  the  sulfate  ion  in  the  solution. 

The  normal  ionic  reaction  is  written 

Cr2O7"  +  6Fe++  +  ^H  ^±1  2Cr+++  +  6Fe+++  +  7H2O. 

We  believe  that  the  reaction  without  acid  proceeds  in  the  same  way, 
the  hydrogen  ion  being  derived  from  the  water. 

H20  ^±1  H+  +  OH". 

As  hydrogen  ion  is  consumed  by  the  reaction,  more  is  formed,  and  at 
the  same  time  hydroxyl  ion  accumulates.  Soon  the  concentration  of 
hydroxyl  ion  is  sufficient  to  exceed  the  solubility  products  of  the  hydroxides 
of  iron  and  chromium,  and  the  colloidal  hydrous  oxides  are  formed. 
Fe+++  +  3QH-  5±  Fe(OH)3;  2Fe(OH),  +  (y  —  3)H2O  ^±  Fe2O3.jH2O. 
Cr+++  +  30H-  ^±  Cr(OH)3;  2Cr(OH)3  +  (*  —  3)H2O  Z£±  Cr2O3.*H2O. 

Summary. 

1.  The   stoichiometric   relations   in    the   reaction   between   potassium 
dichromate  and  ferrous  sulfate  are  the  same  with  or  without  acid. 

2.  The  experimental  conditions  for  the  complete  precipitation  of  ferrous 
iron  by  ammonium  hydroxide  have  been  found  and  employed  to  determine 
dichromate  in  a  mixture  also  containing  ferrous,  ferric,  and  chromium  salts. 

3.  Without  acid  the  reaction  is  instantaneous,  except  in  very  dilute 
solutions. 

4.  Disregarding  the  change  of  hydrogen-ion  concentration  accompanying 
the  reaction,  the  order  is  higher  than  the  fourth. 

5.  The  rate  of  the  reaction,  with  acid,  can  not  be  proportional  to  the 
second  power  of  the  concentration  of  acid  added,  for  then  it  should  be 
zero  without  acid. 

6.  The  products  of  the  reaction  are  the  sulfates  of  potassium,  chromium, 
and  iron,  and  the  colloidal  hydrous  oxides  of  iron  and  chromium.     The 
latter  are  precipitated  by  the  sulfate  ion,  and  adsorb  a  large  quantity 
of  ferric  sulfate  and  smaller  quantities  of  the  other  two  sulfates. 

7.  The  equations  for  the  reaction  have  been  formulated.  - 


OXIDATION  AND  REDUCTION  WITHOUT  THE  ADDITION  OF 

ACID. 

H.    THE   REACTION  BETWEEN    STANNOUS    CHLORIDE   AND    POTASSIUM 
DICHROMATE.     A  CONTRIBUTION  TO  COLLOID-CHEMISTRY. 

It  may  be  stated  from  the  results  of  the  first  paper  of  this  thesis  that 
colloidal  hydrous  oxides  or  hydroxides  are  obtained  in  an  oxidation- 
reduction  reaction,  in  which  acid  must  be  added  for  the  formation  of 
normal  salts,  if  the  stoichiometric  relation  is  the  same  without  acid  as 
with  acid.  If  the  reaction  involves  ions  which  are  good  precipitants 
of  the  colloids  formed,  precipitation  takes  place;  otherwise,  hydrosols 
are  obtained. 

The  equation  for  the  reaction  between  stannous  chloride  and  potassium 
dichromate,  with  acid,  is 

3SnCl2  +  K2Cr207  +  i4HCl  ==  aSnCU  +  2CrCl3  +  yH2O  +  2KC1, 
where,  it  is  seen,  fourteen  mols  of  hydrochloric  acid  per  mol  of  dichromate 
are  necessary  to  form  the  normal  salts  of  tetravalent  tin,  trivalent 
chromium,  and  of  potassium.  The  object  of  this  investigation  was  to 
determine  whether  the  stoichiometric  relation  between  dichromate  and 
stannous  chloride  is  the  same,  i.  e.,  one  mol  of  the  former  oxidizing  three 
mols  of  the  latter,  and  what  substances  are  formed  when  no  acid  is  added. 
The  Stoichiometric  Relation. 

Samples  of  commercial  c.  P.  stannous  chloride  of  about  0.4  g.  each  were 
rapidly  transferred  to  beakers  from  a  weighing  bottle,  dissolved  in  50  cc. 
of  water,  and  the  solutions  titrated  with  standard  dichromate,  some  after 
adding  10  cc.  of  concentrated  hydrochloric  acid  and  others  without  the 


addition  of  any  acid.     Ferrous  ammonium  sulfate  solution  containing 
potassium  thiocyanate  was  employed  as  an  outside  indicator. 

At  first  the  results  of  the  titrations  without  acid  seemed  to  be  slightly 
higher  than  those  with  acid,  which,  as  in  the  titrations  of  ferrous  sulfate 
with  dichr ornate  without  the  addition  of  acid,  would  indicate  that  the 
reaction  was  not  instantaneous.  Further  investigation,  however,  showed 
that  the  hydrochloric  acid  alone  gave  a  faint  pink  color  with  the  indicator, 
which  was  caused  by  the  ferric  iron  in  the  slightly  oxidized  ferrous  sulfate 
of  the  indicator  being  brought  into  solution  by  the  acid.  To  overcome 
this  difficulty,  the  titrations  in  which  acid  was  used  were  run  until  a  drop 
of  the  solution  gave  a  darker  tint  with  the  indicator  than  acid  alone. 
The  results  with  and  without  acid  were  then  almost  identical. 

Therefore  the  oxidizing  power  of  dichromate  towards  stannous  chloride 
is  not  affected  if  the  reaction  takes  place  without  the  addition  of  acid. 
Furthermore,  the  reaction  is  practically  instantaneous  in  dilute  solutions, 
or,  in  the  titrations  referred  to  above,  the  results  should  be  higher  without 
acid  than  with  acid.  It  is  not  surprising,  however,  that  this  should  be 
the  case,  for  the  reaction  may  be  regarded  as  compounded  of  two,  each  of 
which  has  a  very  great  velocity,  namely,  that  between  ferrous  salt  and 
dichromate  and  that  between  ferric  salt  and  stannous  salt. 

When  no  acid  is  used,  stannous  chloride  is  considerably  hydrolyzed 
in  the  concentrations  employed  in  our  titrations,  giving  milky,  opalescent, 
solutions.  The  dichromate  rapidly  reduced  the  turbidity,  which  dis- 
appeared after  a  few  cubic  centimeters  had  been  added.  It  is  also  in- 
teresting to  note  that  at  the  end  of  these  titrations,  the  solutions  possessed 
a  peculiar,  faint  and  yet  distinct,  fruity  odor,  which  was  not  observed 
in  the  acid  titrations.  We  have  been  unable  to  discover  the  cause  of  this 
odor.  In  titrating  the  solid  salt  with  o.i  N  dichromate,  an  olive-green 
solution  having  no  turbidity  is  obtained  immediately. 
The  Products  of  the  Reaction. 

The  percentage  of  stannous  tin  contained  in  the  solid  chloride  was 
estimated  by  titration  with  o.  i  N  dichromate  in  the  presence  of  an  excess 
of  hydrochloric  acid,  and  the  amount  containing  an  equivalent  weight  in 
grams  calculated.  This  quantity,  119.6  g.,  was  dissolved  in  about  300 
cc.  of  water  contained  in  a  liter  flask.  An  equivalent  weight  of  potassium 
dichromate  (49.03  g.),  enough  to  completely  oxidize  the  stannous  chloride, 
was  dissolved  in  200-300  cc.  of  water  contained  in  a  beaker.  The  di- 
chromate solution  was  gradually  added  to  the  stannous  chloride  solution, 
the  mixture  being  shaken  vigorously  to  secure  homogeneity.  During 
the  process  of  mixing,  brownish  and  greenish  blue  gelatinous  masses 
were  formed  and  at  one  point  the  entire  mixture  became  a  jelly;  but 
when  all  the  dichromate  had  been  added  a  perfectly  clear,  deep,  olive- 


15 

green  solution  resulted,  which  in  sufficient  depth  appeared  red  by  trans- 
mitted light,  natural  or  artificial.  The  mixture  was  diluted  to  a  liter 
and  aliquot  portions  removed  for  investigation. 

Treatment  with  Ethyl  Alcohol. — One  hundred  cubic  centimeters  were 
evaporated  to  dryness  on  a  steam  bath  and  dried  to  constant  weight  in  an 
air  oven.  The  residue  was  treated  with  95%  ethyl  alcohol,  which  dis- 
solved all  but  a  white  crystalline  substance  slightly  tinged  with  green. 
This  alcohol-insoluble  matter  was  filtered  off  by  suction  through  a 
Biichner  funnel  and  washed  with  95%  alcohol,  but  it  could  not  be  en- 
tirely freed  from  the  slight  coloration  due  to  an  adsorbed  chromium 
compound. 

Alcohol-Insoluble  Matter. — The  residue  from  100  cc.  of  the  original 
mixture  was  dissolved  in  water  and  made  up  to  250  cc.  Twenty-five 
cubic  centimeter  portions  were  removed  for  the  estimation  of  tin, 
chromium  and  chlorine.  The  tin  was  determined  by  precipitation  with 
hydrogen  sulfide  and  ignition  to  stannic  oxide;  the  chromium  by  addition 
of  ammonium  hydroxide  to  the  hydrogen  sulfide  filtrate;  and  the  chlorine 
by  precipitation  with  silver  nitrate.  The  potassium  was  obtained  by 
subtracting  the  amount  in  the  alcohol-soluble  matter  from  the  total 
employed  in  the  reaction. 

Alcohol-Soluble  Matter. — The  alcohol  solution  was  evaporated  to 
dryness  on  a  steam  bath,  and  the  residue,  dried  to  constant  weight  in  an 
air  oven  at  120°,  ground  and  analyzed  for  potassium,  chlorine,  chromium 
and  tin.  The  methods  for  the  tin  and  chlorine  were  the  same  as  those 
above,  while  the  potassium  was  determined  as  potassium  chloride,  and  the 
chromium  by  fusion  with  sodium  peroxide  and  titrating  the  resulting 
chromate  with  thiosulfate. 

The  results  calculated  to  totals  for  the  entire  mixture  are  as  follows: 

Alcohol-insoluble  matter.  Alcohol-soluble  matter. 


K 


Crm. 

Sniv. 


Grams. 

Gram 
equivalents. 

Grams. 

Gram 
equivalents. 

JLUIH.1 

gram 
equiv. 

12.  II 

0.3097 

0.92 

0.0235 

0.3332 

II.  12 

0.3133 

19.01 

0.5361 

0.8494 

0.15 

0.0087 

17.15 

0.9894 

0.9981 

0-45 

O.OI5I 

60.39 

2.0300 

2.0451 

The  quantities  of  the  elements  contained  in  the  entire  mixture  are: 
potassium,  0.3333  equivalent;  chromium,  i  equivalent;  chlorine  and 
tin,  i  and  2  equivalents,  respectively,  provided  the  stannous  chloride  did 
not  contain  stannic  tin.  It  will  be  remembered  that  the  weight  of  stannous 
salt  containing  one-half  the  molecular  weight  of  unoxidized  chloride  was 
employed.  The  results  show  that  0.0451  equivalent  of  stannic  tin  was 
present,  with  which  0.0226  equivalent  of  chlorine  was  associated.  Thus, 
the  total  chlorine  should  be  1.0226  equivalents. 


i6 

It  is  evident  that  the  substance  separated  in  the  alcohol  treatment  is 
potassium  chloride,  which  therefore  is  one  of  the  products  of  the  reaction. 

The  constituents  of  the  alcohol-soluble  matter  cannot  be  determined 
from  the  analysis  alone.  It  may  be  observed,  however,  that  there  is  a 
deficiency  of  chlorine  of  nearly  one-sixth  the  total.  This  loss  was  in- 
curred when  the  alcohol-soluble  matter  was  dried  in  the  air  oven,  and 
was  due  to  the  decomposition  of  stannic  chloride  or  chromic  chloride,  or 
both. 

Since  no  tin  was  lost  in  the  process  of  drying,  hydrated  stannic  chloride 
could  not  have  been  present  to  any  appreciable  degree,  for  hydrated 
stannic  chloride  volatilizes  considerably  when  heated.1  Hydrated  chromic 
chloride  on  the  other  hand  does  yield  hydrochloric  acid  when  heated  to 
120°  C.2  We  are  thus  led  to  the  conclusion  that  the  alcohol-soluble 
matter  consists  of  a  mixture  of  the  hydrous  oxides  of  tin  and  chromium, 
and  hydrated  chromic  chloride. 

If  stannic  chloride  were  a  product  of  the  reaction,  it  could  be  extracted 
by  means  of  carbon  bisulfide.  A  portion  of  the  original  mixture  was 
evaporated  to  dryness  on  a  water  bath,  the  residue  powdered,  introduced 
into  a  thimble,  and  extracted  with  carbon  bisulfide  for  about  eighteen 
hours,  but  no  trace  of  tin  could  be  found  in  the  solvent.  Dialysis  of  the 
reaction  mixture  gave  further  evidence  that  it  does  not  contain  stannic 

chloride. 

Dialysis. 

Fifty  cubic  centimeters  of  the  original  solution  were  dialyzed  in  a 
parchment  paper  bag  suspended  in  a  beaker  filled  with  distilled  water 
to  the  level  of  the  solution  in  the  bag.  In  a  short  time  the  external  liquid 
was  colored  bluish  green  and  considerable  osmosis  had  taken  place. 
Fresh  water  was  placed  in  the  beaker  every  day  until  it  was  not  per- 
ceptibly colored  after  standing  twenty-four  hours.  The  accumulated 
diffusate  for  this  period  was  concentrated  and  tested  for  tin,  but  none 
was  found,  thus  proving  the  absence  of  stannic  chloride  in  the  mixture. 

The  dialysis  was  continued  in  order  to  free  the  hydrosol  as  far  as  possible 
from  electrolytes.  Excessive  dilution  by  osmosis  was  avoided  by  keeping 
the  level  inside  the  membrane  several  centimeters  higher  than  outside, 
which  resulted  in  the  concentration  of  the  colloid  as  the  osmosis  dimin- 
ished. In  about  five  weeks  when  the  diffusate  was  giving  only  a  faint 
test  for  chloride  ion,  the  contents  of  the  bag  set  to  a  perfectly  clear  gel 
of  a  beautiful,  emerald  green  in  reflected  light,  and  a  deep  red  in  trans- 
mitted light.  In  a  second  dialysis  where  no  effort  was  made  to  keep 
the  solution  in  the  membrane  from  increasing  in  volume,  250  cc.  in- 
creased to  1 1 oo  cc.  in  five  weeks,  giving  a  clear  hydrosol  resembling  the 

1  Gmelin-Kraut,  4,  I,  313. 

1  Itnd.t  3, 1,  439- 


17 

hydrogel.  The  hydrosol  may  be  boiled  down  to  a  very  viscous  con- 
sistency and  on  being  dehydrated  over  sulfuric  acid  becomes  a  firm  gel. 

Analysis  of  the  gel  showed  that  it  contained  all  the  tin,  practically  one- 
half  of  the  chromium  and  a  negligible  quantity  of  chlorine.     These  re- 
sults, together  with  those  previously  obtained,   enable  us  to  formulate 
the  reaction  as  follows : 
2K2Cr207  +  6SnCl2  +  (6*  +  ?)H2O  ^± 

4KC1  +  6SnO2.#H2O  +  Cr2O3.;yH2O  +  2CrCU  +  2HC1 

Thus  written  the  equation  expresses  the  fact  that  dialysis  yields  a 
mixed  hydrosol  of  hydrous  stannic  and  chromic  oxides  in  the  molecular 
ration  of  6  SnO2  to  i  Cr2O3. 

The  reaction,  with  acid,  is  written  ionically 

Cr207-  +  3811++  +  HH+  Z^±  2Cr+++  +  3Sn+++++  7H2O.    (i) 

Though  no  acid  is  added  in  our  reaction,  hydrogen  ion  is  present,  due 
to  the  hydrolytic  dissociation  of  the  dichromate  and  stannous  chloride. 
Equation  i,  therefore,  represents  the  reaction  when  no  acid  is  added. 

Aqueous  solutions  contain  hydroxyl  ion  in  equilibrium  with  hydrogen 
ion  according  to  the  equation: 

H+  +  OH-  ^±  H20  (2) 

As  the  hydrogen  ion  is  removed  by  reaction  (i)  this  equilibrium  is  dis- 
turbed, the  more  so  as  the  available  hydrogen  is  very  limited  and  the 
concentration  of  hydroxyl  ion  correspondingly  increases.  When  the 
latter  accumulates  sufficiently,  the  solubility  products  of  stannic  and 
chromic  hydroxides  are  exceeded  and  the  colloidal  hydrous  oxides  are 
formed 

3OH-  ^±  Cr(OH)3;  2Cr(OH)3  +  (y  —  3)H2O  ^±  Cr2O3.;yH2O. 
4OH-  ^±  Sn(OH)4;  Sn(OH)4  +  (x  —  2)H2O  ^±SnO2.^H2O . 
The  formation  of  the  hydrosols  maintains  the  concentration  of  hydroxyl 
ion  within  perfectly  definite  limits,  which  insures  a  definite  minimal 
concentration  of  hydrogen  ion  sufficient  for  reaction  (i),  which  there- 
fore continues  to  completion. 

Conclusions. 

1.  The   stoichiometric   relations   in    the   reaction    between   potassium 
dichromate  and  stannous  chloride  are  the  same  with  or  without  acid. 

2.  The  products  of  the  reaction  are  potassium  and  chromium  chlorides, 
and  stannic  and  chromic  hydrous  oxides  in  colloidal  solution. 

3.  A   clear  mixed  hydrosol   of  stannic  and   chromic  hydrous  oxides, 
approximately   in   the   molar   ratio   of   6  SnO2  to  i   Cr2O3,  may   be   ob- 
tained by  adding  an  equivalent  amount  of  dichromate  solution  to  stan- 
nous chloride    and  dialyzing  the    mixture.     The    hydrosol   will  contain 
all  of  the  tin  and  practically  one-half  of  the  chromium  used  in  the  re- 
action. 

4.  The  equations  for  the  reaction  have  been  formulated. 


VITA. 

Joshua  Chitwood  Witt  was  born  at  Connersville,  Indiana,  1884.  He 
attended  school  at  Connersville  and  Liberty,  Indiana,  and  was  graduated 
from  high  school  in  1903.  He  entered  Butler  College,  Indianapolis,  in 
1904,  receiving  the  degree  of  Bachelor  of  Arts  in  1908,  with  chemistry  as 
his  major  subject.  The  same  year  he  took  up  the  study  of  chemistry  and 
bacteriology  at  the  University  of  Chicago,  which  culminated  in  the  degree 
of  Bachelor  of  Science.  From  1909  to  1911,  he  did  special  work  in 
mechanical  engineering  at  Armour  Institute  of  Technology.  In  1911  he 
entered  the  University  of  Pittsburgh  to  pursue  graduate  work  in  chemistry 
and  physics,  receiving  the  degree  of  Master  of  Science  the  following  year. 


BATB 


AN  INITIAL  FINE  OF  25  CENTS 

OVERDUE. 


W6~»8  1993 


LD  21-100m-8,'34 


Photomount 
Pamphlet 

Binder 
Gaylord  Bros. 

Makers 
Syracuse,  N.  Y. 

PAT.  JAN  21,  1908 


U.C.  BERKELEY  LIBRAR 


m 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


